C32 E07 Remote Hot Cold Packs PDF

Preview

Summary

Lab assignments ...

Description

CHEM 032 Experiment 7: Thermodynamics of Hot and Cold Packs

Lab Report Rubric

Max Points

Purpose: Purpose

1.0

Introduction: Background Information: Entropy Gibbs Free Energy • Standard Conditions Procedure & Techniques: • •

Procedure Summary

3.0

2.0

Equipment & Techniques: • Thermometers • Measuring Temperature Safety & Waste:

2.0

Safety Considerations

1.0

Waste Disposal

1.0

Data & Calculations: •

Calculations

4.0

•

Graphs

2.0

•

Post-lab questions

1.0

Results & Discussion: •

Reports main results

2.0

•

Explains results conceptually

2.0

•

Error analysis

2.0

Presentation Points: •

All required elements are present

•

Material is organized

2.0

Lab Report Total:

25.0

Your Points

CHEM 032 Experiment 7: Thermodynamics of Hot and Cold Packs

Lab Reports: The laboratory report will consist of the parts listed below and summarized on the rubric. The laboratory report should be submitted as a single file (.pdf is preferred, but .doc is also acceptable) through the Blackboard assignment tool. Laboratory reports are due 1 week after the assignment has been opened. If handed in late, a 3 point deduction will be made for every day a report is late. No lab reports will be accepted that are more than one week (7 days) late. Purpose: In a sentence or two, summarize the purpose of the lab or the questions to be answered during the lab. The purpose should be concise and specific to the goal of the particular experiment. Note that in some cases an experiment may have more than one objective and that you should include them all in the purpose statement. Introduction: The introduction should not only define the key background information listed, but also relate these key points back to how they specifically apply to the experiment being discussed. Procedure & Techniques: The procedure summary should provide a brief overview in paragraph form of the procedure that was used. Using the lab handouts and videos, describe in paragraph form the experimental procedure used. New techniques or equipment used should be explained. If the equipment or technique has been used before, explain why and what it is being used for again. Safety & Waste: The safety and waste section should include all hazards present in the experiment, precautions to minimize the hazards, and proper waste disposal for the experiment. Data & Calculations: The calculation and questions provided on the data sheets must be competed and included as a part of the lab report. You should not type out equations and calculations; instead it is suggested that you complete these by hand and included a scanned or picture version as part of your complete lab report. Results & Discussion: Review all the data from the experiment. In paragraph form, summarize the main finding(s) of this lab. The main experimental result(s) should be clearly stated and related back to the purpose of the experiment. At least two potential experimental errors, their possible effects, and ways to reduce errors should also be discussed. Presentation Points: The lab report is typed and uses headings and subheadings to visually organize the material. All required elements are present. Few errors in spelling, punctuation and grammar in the report.

CHEM 032 Experiment 7: Thermodynamics of Hot and Cold Packs Background Prior to writing up your pre-lab and attending your laboratory section you should review the following sections from the course textbook in preparation for the experiment: Review Chapter 9 – Thermochemistry 9.4 Quantifying Heat and Work 9.5 Measuring ∆𝐸 for Chemical Reactions: Constant-Volume Calorimetry 9.6 Enthalpy: The Heat Evolved in a Chemical Reaction at Constant Pressure 9.7 Measuring ∆𝐻 for Chemical Reactions: Constant-Pressure Calorimetry Chapter 18 – Free Energy and Thermodynamics 18.1 Nature’s Heat Tax: You Can’t Win and You Can’t Break Even 18.2 Spontaneous and Nonspontaneous Processes 18.3 Entropy and the Second Law of Thermodynamics 18.4 Entropy Changes Associated with State Changes 18.5 Heat Transfer and Entropy Changes of the Surroundings 18.6 Gibbs Free Energy ° 18.7 Entropy Changes in Chemical Reaction: Calculating ∆𝑆%&' ° 18.8 Free Energy Changes in Chemical Reactions: Calculating ∆𝐺%&'

Introduction There are a number of small portable hot or cold packs on the market for the rapid treatment of injuries, storage of delicate medical supplies or protection from extreme cold. All of these are designed for use away from traditional power supplies (electricity) and therefore make direct use of chemical reactions in order to either release or absorb thermal energy accordingly. A typical hot or cold pack consist of a plastic bag with two compartments, one containing water and the other a salt. When the bag is broken, the contents of the two compartments mix, dissolving the salt. How hot or cold the pack gets depends on the salt’s concentration in the water. Hot packs can reach temperatures as high as 90ºC. Cold packs can reach temperatures as low as 0ºC. Instant hot and cold packs are frequently used by athletes to treat injuries. For example, cold packs are routinely applied to sprained ankles. Heat flows from the ankle to the pack, reducing the temperature of the injured area. Lowering the injured ankle’s temperature produces vasoconstriction of the blood vessels that reduces blood flow, which in turn reduces inflammation. Hot packs can be applied to reduce muscle spasms, muscle soreness, inflammation, and relieve pain. Heat flows from the pack to the affected area, increasing its temperature. This produces vasodilation, which increases blood flow into the target tissue. Increased blood flow brings needed oxygen and nutrients to the injured area, aiding the healing process.

E07-1

CHEM 032 Experiment 7: Thermodynamics of Hot and Cold Packs When ionic compounds dissolve in water, they either absorb energy from or release energy to the surroundings. A chemical reaction that absorbs heat from the surroundings is an endothermic process. A chemical reaction that releases heat to the surroundings is an exothermic process. The energy or enthalpy change associated with the process of a solute dissolving in a solvent is called the heat of solution (∆H+,-. ). In the case of an ionic compound dissolving in water, the overall energy change is the net result of two processes: the energy required to break the attractive forces (ionic bonds) between the ions in the crystal lattice, and the energy released when the dissociated (free) ions form ion-dipole attractive forces with the water molecules. Heats of solution and other enthalpy changes are generally measured in an insulated vessel called a calorimeter that reduces or prevents heat loss to the atmosphere outside the reaction vessel. The coffee-cup calorimeter consists of two Styrofoam coffee cups, one inserted into the other, to provide insulation from the laboratory environment. The calorimeter is equipped with a thermometer and a stirrer. The reaction occurs in a specifically measured quantity of solution within the calorimeter, so that the mass of solution is known. During the reaction, the heat evolved (or absorbed) causes a temperature change in the solution, which the thermometer measures directly. The process of a solute dissolving in water may either release heat into the aqueous solution or absorb heat from the solution, but the amount of heat exchange between the calorimeter and the outside surroundings should be minimal. The enthalpy changes of dissolution for a salt will determine whether a particular salt is appropriate as a hot or cold pack ingredient. By definition, exothermic processes release heat. That heat can then increase the temperature of the surroundings. If the surroundings include a strained or sore muscle, you have a useful hot pack. For a cold pack, an endothermic reaction that absorbs heat from the surroundings is required. Heat will flow from the surroundings, for example a twisted ankle, and its temperature will be reduced. A non-trivial part of selecting a salt is determining just how much solid should be used in your calorimeter to produce the desired temperature change. qsoln 3 = 3m soln × Cs,soln × ∆T+,-.

(from Section 9.7)

Where:3m is the mass of the reaction solution, C is the specific heat of the solution (for water, 4.18 J/g°C), and ∆T is the temperature change of the solution: (∆T = Tfinal - Tinitial ). The mass of the solutions and the temperature changes of the solutions can and will be measured experimentally. Also recall that: qrxn = −3qsoln

E07-2

CHEM 032 Experiment 7: Thermodynamics of Hot and Cold Packs In the real world it is impossible to construct the perfect calorimeter, that is one in which there is no heat loss or gain except for the reaction being studied. Because of this, measuring the temperature change is complicated by heat leaking into or out of the calorimeter. To correct for these effects a graphical procedure is used to obtain a more accurate value of ΔT. A typical time versus temperature plot for a calorimetric experiment involving an exothermic reaction follows. The initial temperature of the reaction vessel in the experiment represented by this plot is a little below room temperature so heat is initially leaking in. This explains the slow rise in temperature before the reaction is initiated. On mixing the two solutions there is an initial sharp rise in temperature, caused by heat released from the reaction, after which the temperature begins to decrease, a result of heat leaking out of the calorimeter. After a time the reaction is complete and the temperature continues to decrease in a linear fashion. Extrapolation of this linear portion of the curve back to the time the reaction began, when the initial mixing took place, gives a fairly accurate value for what temperature would have been obtained if the reaction and equilibrium had taken place instantaneously, or the final temperature. With this, one can more accurately calculate the change in temperature for the reaction. Since the amount of heat released or absorbed from a specific chemical reaction is dependent upon the size of the chemical reaction carried out (i.e. the number of moles of limiting reagent used) it is customary to correct the observed heat flow for the reaction size by dividing q;...